Classification of Matter

Classification of Matter

• Properties are characteristics of matter used to categorize different types of matter.
• Matter can be categorized by its state and its composition.

Three States of Matter

• Gas state involves particles widely separated with no definite shape or volume.
• Liquid state involves particles closer together with a definite volume but no definite shape.
• Solid state involves particles very close together with a definite shape and volume.

Composition of Matter

• Pure Substance has only one component.
• Mixture is a combination of two or more pure substances where each retains its identity and does not undergo a chemical reaction.
• Element pure substance that cannot be changed into a simpler form by any chemical reaction (e.g., Oxygen, Hydrogen).
• Compound is a pure substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio (e.g., Salt, Water).
• Homogeneous Mixture features uniform composition with well-mixed particles (e.g., Air, Ethanol in Water).
• Heterogeneous Mixture features non-uniform composition with random placement (e.g., Oil and Water, Salt and Pepper).

Physical and Chemical Properties

• Physical Property is observed without changing the composition or identity of a substance.
• Physical Change produces a recognizable difference in the appearance of a substance without causing any change in its composition or identity, like conversion from one physical state to another or melting.
• Chemical Property results in a change in composition and can be observed only through a chemical reaction.
• Chemical Reaction (chemical change) converts a chemical substance into one or more different substances by rearranging, removing, replacing, or adding atoms.

Intensive and Extensive Properties

• Intensive Properties of matter that are independent of the quantity of the substance for example color and melting point.
• Extensive Properties of matter depends on the quantity of the substance for example mass and volume.

Units of Measurement

• Units are the basic quantity of mass, volume, or any quantity being measured.
• A measurement is useless without its units.
• The English system is a collection of functionally unrelated units that are difficult to convert from one to another.
• 1 foot = 12 inches = 0.33 yard = 1/5280 miles in the English system.
• The Metric System consists of units related to each other decimally, making it systematic.
• Metric units relate by powers of tens.
• Mass is the quantity of matter in an object, and is not synonymous with weight.
• Weight = mass × acceleration due to gravity.
• The standard unit of mass is the gram (g).
• 1 lb = 453.6 g.
• Mass must be measured on a balance, not a scale.
• Length is the distance between two points.
• The standard unit of length is the meter (m).
• 1 yd = 0.9144 m.
• Volume is the space occupied by an object.
• The standard unit of volume is the liter (L).
• The metric unit for time is the second (s).
• 1 qt = 0.9464 L

Metric System Prefixes

• Mega (M): 10^6, 1,000,000, 1 Mx = 10^6 x
• Kilo (k): 10^3, 1,000, 1 kx = 10^3 x
• Deka (da): 10^1, 10, 1 dax = 10 x
• Deci (d): 10^-1, 0.1, 1 dx = 10^-1 x
• Centi (c): 10^-2, 0.01, 1 cx = 10^-2 x
• Milli (m): 10^-3, 0.001, 1 mx = 10^-3 x
• Micro (μ): 10^-6, 0.000001, 1 μx = 10^-6 x
• Nano (n): 10^-9, 0.000000001, 1 nx = 10^-9 x

Conversions

• Length × width × height = Volume
• 1 dm × 1 dm × 1 dm = 1 dm^3 = Volume
• 1 dm^3 =1 L

Measurement

• Information-bearing digits or figures in a number are significant figures.
• A measuring device determines the number of significant figures in a measurement.
• The degree of uncertainty in a measurement is indicated by the number of figures used to represent the information.
• Significant figures consist of all digits in a number representing data or results that are known with certainty plus one uncertain digit.
• All nonzero digits are significant, for example 7.314 has four significant digits.
• The number of significant digits is independent of the position of the decimal point.
• Zeros located between nonzero digits are significant, for example 60.052 has five significant digits.
• Zeros at the end of a number (trailing zeros) are significant if the number contains a decimal point, for example 4.70 has three significant digits.
• Zeros that are insignificant if the number does not contain a decimal point, for example 100 has one significant digit but 100. has three.
• Zeros to the left of the first nonzero integer are not significant for example 0.0032 has two significant digits.
• Scientific Notation is used to express very large or very small numbers easily and with the correct number of significant figures.
• It represents a number as a power of ten.
• To convert a number greater than 1 to scientific notation, the original decimal point is moved x places to the left, and the resulting number is multiplied by 10^x.
• The exponent x is a positive number equal to the number of places the decimal point moved.
• To convert a number less than 1 to scientific notation, the original decimal point is moved x places to the right, and the resulting number is multiplied by 10^-x.
• The exponent x is a negative number equal to the number of places the decimal point moved.
• Scientific notation should be used to input very large or small numbers into a calculator( e.g., 0.0000000000000000000000066466 g must be entered into calculator as 6.6466 ×10^-24).
• Accuracy is the degree of agreement between the true value and the measured value.
• Error is the difference between the true value and our estimation, and can be random or systematic.
• Precision is a measure of the agreement of replicate measurements.
• Deviation is amount of variation present in a set of replicate measurements.
• Inexact numbers have uncertainty (degree of doubt in final significant digit).
• Exact numbers are a consequence of counting.
• A set of counted items (beakers on a shelf) has no uncertainty.
• Exact numbers by definition have an infinite number of significant figures.
• When the number to be dropped is less than 5, the preceding number is not changed.
• When the number to be dropped is 5 or larger, the preceding number is increased by one unit.
• The result in a calculation cannot have greater significance than any of the quantities that produced the result.
• When adding and subtracting numbers in scientific notation, the numbers must have the same exponent

Unit Conversion

• The Factor-Label Method (Dimensional Analysis) uses conversion factors to convert from one unit to another within the same system and to convert units from one system to another.
• To convert from one unit to another you must know the conversion factor, which is the relationship between the two units.
• Example relationship, 1 gal = 4 qt.
• Example conversion factor, 1 gal / 4 qt.
• To use conversion factors, write the data given, then multiply by the conversion factor with the unit of the data given in the denominator to produce the desired result.
• To multistep convert, use a data of conversion factors with the unit of the data given in the denominator in each step to convert to the final desired unit

Quantities

• Temperature- the degree of "hotness" of an object
• Kelvin (K) scale is another temperature scale. It is of particular importance because it is directly related to molecular motion.
• As molecular speed increases, the Kelvin temperature proportionately increases.
• Tk =TC +273.15
• Energy is the ability to do work, with forms including kinetic energy (energy of motion) and potential energy (stored energy).
• Energy categorized by form including light, heat, electrical, mechanical and chemical.
• Energy cannot be created or destroyed.
• Energy may be converted from one form to another.
• Energy conversion always occurs with less than 100% efficiency.
• All chemical reactions involve either a "gain" or "loss" of energy.
• Basic Units of Energy including
• calorie or joule,
• 1 calorie (cal) = 4.184 joules (J).
• kilocalorie (kcal) = food Calorie.
• 1 kcal = 1 Calorie = 1000 calories
• A calorie is the amount of heat energy required to increase the temperature of 1 gram of water 1 degree C.
• Concentration is the number or mass of particles of a substance contained in a specified volume.
• Concentration is often used to represent various mixtures of different substances, for example concentration of oxygen in the air, pollen counts, or the proper dose of an antibiotic.
• Density is the ratio of mass to volume.
• Density (d) = mass / volume = m / V.
• Density is an extensive property.
• Density is used to characterize a substance, each substance has a unique density.
• Density may be measured as g/mL or g/cm^3 or g/cc.
• Specific Gravity is a unit less values of density are often related to a standard.
• Specific gravity is the ratio of the density of the object in question to the density of pure water at 4 degrees C.
• The unit less specific gravity means the 2 units cancel.
• specific gravity = (density of object (g/mL)) / (density of water (g/mL))
• Health industry uses specific gravity to test urine and blood samples.